Question

Make a sketch of an electrochemical cell with the overall reaction shown here. Label the anode, the cathode, and the salt bridge. Indicate the direction of electron flow. Hint: When drawing electrochemical cells, the anode is usually drawn on the left side. $$ \mathrm{Mg}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{Mg}^{2+}(a q)+\mathrm{Ni}(s) $$

Short answer

On the left of the sketch is the anode, where \( \mathrm{Mg} \longrightarrow \mathrm{Mg}^{2+} + 2e^- \) occurs. On the right is the cathode where \( \mathrm{Ni}^{2+} + 2e^- \longrightarrow \mathrm{Ni} \) happens. The salt bridge connects the two compartments, and electrons flow from the anode to the cathode through an external wire.

Step-by-step solution

Identify Anode and Cathode Reactions

Determine the oxidation and reduction half-reactions. The metal that gets oxidized (loses electrons) is the anode and the metal that gets reduced (gains electrons) is the cathode. For the given reaction, \( \mathrm{Mg}(s) \longrightarrow \mathrm{Mg}^{2+}(aq) + 2e^- \) is the oxidation half-reaction, so magnesium is the anode. The reduction half-reaction is \( \mathrm{Ni}^{2+}(aq) + 2e^- \longrightarrow \mathrm{Ni}(s) \) which means nickel is the cathode.

Draw Electrochemical Cell

On the left side of your sketch, draw the anode compartment with \( \mathrm{Mg}(s) \) and label it 'Anode'. Sketch a wire connected to the anode leading to an external circuit. On the right side, draw the cathode compartment with \( \mathrm{Ni}^{2+}(aq) \) and label it 'Cathode'. Connect this also with a wire to the external circuit. To complete the circuit, ensure the wires meet at a point which could represent a voltmeter or other device to complete the circuit.

Include the Salt Bridge

Draw a salt bridge, usually represented as an inverted U-tube that connects the two compartments. The salt bridge contains an electrolyte that permits the flow of ions, but not the mixing of different solutions, which would allow the electrical neutrality to be maintained in each compartment.

Indicate Electron Flow

Indicate the direction of electron flow from the anode to the cathode in the external circuit. Electrons flow from where they are lost (anode) to where they are gained (cathode). A standard convention is showing electron flow with a half-arrow pointing in the direction of the flow along the external wire.

Label the Electrolyte Solutions

Label the anode compartment as containing \( \mathrm{Mg}^{2+}(aq) \) ions, and the cathode compartment as containing \( \mathrm{Ni}^{2+}(aq) \) ions. This indicates the products of the electrode reactions and substances in the electrolyte.

Key concepts

Anode and Cathode Reactions

Understanding the reactions at the anode and cathode is essential when studying electrochemical cells. These reactions are the heartbeat of the cell, driving the chemical process that generates electrical energy.

At the anode, oxidation occurs. This means that a substance loses electrons and is therefore referred to as the oxidation half-reaction. For instance, in our exercise where magnesium is the substance at the anode, the reaction is as follows: \( \text{Mg}(s) \rightarrow \text{Mg}^{2+}(aq) + 2e^- \). Here, solid magnesium (\text{Mg}) loses two electrons (2e^-) and becomes magnesium ions \( \text{Mg}^{2+} \).

Conversely, at the cathode, reduction takes place; a substance gains electrons, as showcased in the reduction half-reaction. In the given cell, nickel ions (\text{Ni}^{2+}) gain two electrons to form solid nickel (\text{Ni}), represented by the equation \( \text{Ni}^{2+}(aq) + 2e^- \rightarrow \text{Ni}(s) \).

Through these reactions, the function of each electrode is defined: the anode is where oxidation occurs, and the cathode is where reduction happens. As students improve their understanding, visualizing these processes can be helpful, allowing them to predict the behavior of other metals in similar setups.

Electron Flow in Electrochemistry

Electron flow in electrochemistry is akin to water flowing down a hill; it's all about moving from a high energy level to a lower one. In an electrochemical cell, electrons flow through an external circuit from the anode to the cathode during a redox reaction.

Why does this happen? It's because the anode is where oxidation (loss of electrons) occurs, creating an excess of electrons. The cathode, on the other hand, is where reduction (gain of electrons) happens, providing a 'sink' for the electrons. The flow is driven by a difference in potential (or charge) between the two electrodes.

Moreover, this flow of electrons generates electric current, which can be harnessed to do work, such as powering a light bulb or charging a battery. The flow is typically indicated by arrows in diagrams, starting from the anode, passing through the wire (often through a load such as a voltmeter), and ending at the cathode. If students grasp this concept, they can apply it broadly to understand various electrochemical processes and technologies, from batteries to corrosion prevention.

Salt Bridge Function

The salt bridge is an often understated but crucial component of the electrochemical cell, acting as a bridge for ions, not unlike a highway overpass allows cars to move between two points without interference.

Its primary role is to maintain electrical neutrality within the separate compartments of the cell, allowing the cell to continue operating without interruption. It does this by permitting the flow of ions, but not allowing the two different solutions to mix, which could neutralize the cell's potential and stop the reaction.

For example, in our exercise, as the magnesium anode releases electrons, it also releases Mg2+ ions into the solution. Simultaneously, as Ni2+ ions at the cathode accept electrons to become solid Ni, there is a net movement of negative charge through the wire. The salt bridge allows anions to flow into the anode compartment and cations into the cathode compartment to balance the charge due to the electron flow.

Explaining the function of the salt bridge enhances students' understanding of the operation of the electrochemical cell and empowers them to design and analyze various electrochemical systems, including batteries and electrolysis setups.